1. Light: Wave or Particle Chapter 4, Section 1 notes

2. Historically……  Light was considered to be a wave until the year 1900. Why?  It is a type of electromagnetic radiation (see electromagnetic spectrum)  Light moves through air with a constant speed like other forms of e.m. radiation.  Light moves like a wave with a measurable wavelength and frequency.

3. Then…..  After the year 1900, scientists conducted experiments that documented the Photoelectric Effect.  This changed the idea of light because, under the wave theory of light, any frequency of light should have had enough energy to knock loose an electron. But, after conducting experiments, scientists found that only light above a certain frequency would have the necessary energy.

4. Light as a Particle  Max Planck suggested that thermally hot objects emit light in small amounts, called quanta, rather than in a continuous wave of energy.  Technically, a quantum of energy is the minimum amount of energy that an atom can gain or lose.

5. Dual Wave-Particle Nature  After Planck’s theory, Einstein introduced the idea that light isn’t just a wave or a particle, it’s both!  He detailed that a wave of light could also be thought of as a stream of particles in which each particle contained a quantum of energy. Einstein’s term for these particles was a photon (which has zero mass).

6. Back to the Photoelectric Effect  Einstein explained the mystery of the Photoelectric Effect by saying that e.m. radiation is emitted only in whole- numbers of photons.  So, enough energy must hit the object to add up to 1 photon in order to emit an electron. According to Planck’s equation, this means that there must be a minimum frequency, since energy is based on frequency.  This minimum frequency changes for different metals.

7. States of Energy  Atoms exist at different energy levels. There is a ground state (lowest energy level), and an excited state (higher energy level).  Emission is when an electron falls to a lower energy level.  Absorption is when an electron rises to a higher energy level.

8. Emission-Line Spectrum  An Emission-Line Spectrum is when a beam of light shines through a prism, which separates that light into 4 specific colors of the visible spectrum.  These 4 specific wavelengths of color can help identify different atoms and their energy levels.

9. Quantum Theory  Remember, a quantum is the minimum quantity of energy that can be lost or gained by an electron.  The discovery of emission-line spectrum led to the development of quantum theory, based on the emission of a photon from a single Hydrogen atom when the excited H atom falls to its ground state.

10. Bohr and Quantum Theory  Because the emission lines are fixed, then the electrons must exist only in fixed energy states. These fixed energy states led to Bohr’s model of the atom, in which electron are place in certain energy orbitals based on the wavelengths seen in the emission-line spectrum of hydrogen.

11. Shortcomings  All the experimental evidence of the time (1913) led to Bohr’s model of the atom. However, this model works only for Hydrogen, which has only 1 electron. Bohr’s model also doesn’t explain the chemical behavior of some elements (such as valence electrons and bonding).